1. Atomic Structure - Lecture notes 1 PDF

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Lecture on atomic structure ...

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Elements • All matter is composed of elements. • Elements are pure substances that cannot be broken down into simpler substances. • There are 118 elements, only 88 of which occur naturally. • Elements are represented by 1 or 2 letter symbols eg I = iodine, F = fluorine, Fe = iron, Sn =tin • All elements are represented on the Periodic Table • An element’s position on the Periodic Table is very ordered • All elements are composed of atoms • Elements differ from one another due to differences in the structure of the atoms from which they are made

Atoms Major sub-atomic particles are • Electrons – carry a negative charge (one unit of –ve charge) • Protons – carry a positive charge (one unit of +ve charge) • Neutrons – no charge 1. 2. 3.

Mass of an electron: = 9.07 x 10-28 grams Mass of a proton: grams = 1.67 x 10-24 grams Mass of a neutron: 1.67 x 10-24 grams

Mass and Charge of Sub-atomic Particles

Atomic Number and Mass Number • Atomic number (Z) – the number of protons in the nucleus of an atom. • Mass number (A) – the number of protons + number of neutrons in the nucleus of an atom. • For neutral atoms the number of electrons will equal the number of protons. When the number of electrons is different from the number of protons, we have ions (charged atoms). • Atomic number and mass number is represented symbolically as shown below mass number

Atomic Number • The number of protons in the nucleus of an atom is called the ATOMIC NUMBER. It is represented by the symbol Z. • All atoms of a specific element have the same atomic number. • It is specified in the periodic table, just above the symbol of the element. E.g. Lithium has an atomic number of 3 (3 protons in the nucleus) and 3 electrons orbiting the nucleus. ! Mass number • The MASS NUMBER is the sum of protons plus neutrons in the nucleus. It is represented by the symbol A. • Mass numbers are not usually shown on the periodic table

The mass of an electron is insignificant compared to that of a proton. The important thing is the charge Structure of the atom • Protons and neutrons are found in the nucleus of the atom • Since these have the greatest mass, most of the mass of an atom is in the nucleus • Electrons are located in a region of space around the nucleus called the electron cloud • Most of the volume of an atom is due to the electron cloud

Atoms and Elements • Atoms are the smallest indivisible particles of matter (pure elements) that retain the properties of those elements. • Further division into sub-atomic particles destroys the properties of the elements. o e.g. A gold atom is the smallest particle of gold that when further subdivided into sub-atomic particles (protons, electrons and neutrons) loses the characteristics of gold. • The atoms of all elements are made of the same fundamental particles - electrons, protons, neutrons and others (over 100). • It is the arrangements of these sub-atomic particles that makes the atoms different for different elements.

Understanding the Periodic Table • The periodic table has columns (groups) and rows (periods). Columns or groups contain elements with similar properties. • Periods relate to the repetition of the same pattern. • Elements are arranged in terms of increasing atomic numbers. • Periodic law states that when elements are arranged in the order of increasing atomic numbers, elements with similar chemical properties will occur at regular or periodic intervals. • Each element belongs to a period and group in the periodic table. Electronic configuration in atoms • In simple terms, an atom consists of a nucleus and one or more electrons. • Electrons go around the nucleus in specific orbits called principle energy levels or shells. • Electrons can move only in any one of these specific orbits depending upon their energy levels. They cannot randomly go around the nucleus in any orbit. • The electron(s) change orbits by absorbing or releasing energy. • When external energy is added to an electron, the electron absorbs energy and jumps ! to a higher shell (farther from the nucleus). • When the energy source is removed, the electron falls back to the natural state, releasing the excess energy in the form of light or other radiation.

Shells, sub-shells and orbitals • The shells in which the electrons move are designated 1, 2, 3, 4 and so on, 1 being closest to nucleus and increasing shell number being further and further away • Higher shell number means higher energies for the electrons and larger distance between the electron and the nucleus. • The filling order for the first 20 electrons for any atom are:

Atomic Radius The size (or atomic radius) of an atom or ion is dependent on; • the number of electrons in the atom • the shells in which these electrons are located • the charge on the nucleus of the atom

1st shell: 2 electrons 2nd shell: 8 electrons 3rd shell: 8 electrons 4th shell: 18 electrons

• Atomic radii decrease across a period due to an increase in the number of protons attracting the electrons (core charge)

Electron configuration of the first 20 elements 1H 11Na 2He 12Mg 3Li 13Al 4Be 14Si 5B 15P 6C 16S 7N 17Cl 8O 18Ar 9F 19K 10Ne 20Ca

• Atomic radii increase down a group because the electrons are in shells further away from the nucleus

Valence Electrons • The electrons in the outermost shell are called valence electrons. • The number of valence electrons varies from 1 to 8. • These valence electrons participate in the chemical reactions forming bonds. • Atoms react with other atoms in order to get a full outer shell. • This can be achieved by either losing electrons, gaining electrons or sharing electrons during the reaction. • A shell becomes complete when there are eight electrons in it (octet rule) or 2 for the first shell (duet rule). • Elements with a full outermost shell are said to be inert and do not participate in any chemical reaction. • Noble gases He, Ne, A, Kr, Xe etc (last column/group of the periodic table) are inert gases. They all have full outermost shells.

Ionisation Energy First ionisation energy is defined as the energy required to remove the most loosely held electron from an atom in the gaseous state

Ions • Ions are electrically charged species formed when atoms lose or gain electrons in order to get a full valence shell • Positive ions (cations) are formed when atoms lose electrons in order to get a full valence shell

• Ionisation energies tend to increase across a period because the electrons are attracted to an increasingly positively charged nucleus • Ionisation energies decrease down a group because the outermost electrons are in shells that are further away from the nucleus • Ionisation energy is highest for the noble gases because they have full outer shells which gives the atom great stability

• Negative ions (anions) are formed when atoms gain electrons in order to get a full valence shell

Electronegativity Electronegativity is defined as the electron attracting power of an atom. • Electronegativity increases across a period because there are more protons in the nucleus attracting the electrons and the electrons are closer to the nucleus because the size of the atom is decreasing. • Electronegativity decreases down a group because the outermost electrons are getting further away from the positively charged nucleus

Stable vs Unstable • Most of the atoms that make up the world around us are stable. • There are also naturally occurring isotopes that are unstable.

Summary

• Stability and instability is due to the balance or imbalance between the strong repulsive force resulting from the repulsion between positively charge protons in the nucleus and the attractive force between protons and neutrons called the “strong nuclear force” • Unstable nuclei may spontaneously lose energy by emitting a particle and in doing so change into a different element or isotope • Radioisotopes are isotopes with unstable nuclei. These nuclei decay in order to become stable and emit radiation during the process.

• Atoms are made of protons, neutrons and electrons • Protons and neutrons are found in the nucleus and are called nucleons • The properties of these sub atomic particles are;

• The nuclei of atoms (nuclides) are represented using A X where X = symbol for the element Z A = mass number (number of protons + neutrons) Z = atomic number (number of protons) Isotopes • Isotopes are atoms that have the same atomic number (same number of protons, meaning same element) but have different mass numbers (meaning different number of neutrons) in the nuclei. • Eg Hydrogen •There are 3 naturally occurring isotopes of hydrogen It is possible to add neutrons to the nucleus and produce different isotopes. • Common isotopes of hydrogen are; 1. Protium: (1 proton and no neutrons) 2. Deuterium: (1 proton + 1 neutron) 3. Tritium: (1 proton + 2 neutrons) • Most elements are made up of a mixture of isotopes.

From this graph of stable isotopes and radioisotopes, it is evident that for larger nuclei there is a distinct imbalance between the number of protons and neutrons. The ‘line of stability’ of the stable nuclides can be seen as a line that curves away from the N = Z line. Nuclear Radiation • Radiation emitted from the nucleus of an atom • Results from unstable isotopes emitting radiation in order to become stable • Three naturally occurring types of radiation 1. Alpha - α 2. Beta - β 3. Gamma - γ 1. Alpha decay • An unstable nuclei ejects an alpha particle that consists of 2 protons and 2 neutrons (the same as a helium nucleus) • the remaining nucleus is called a “daughter nucleus” and is more stable than the original nucleus

• the process is represented by a decay equation

• The decay equation for the γ decay of the metastable technetium-99 is as follows

Summary • α particles are slow and heavy compared to other radiation types and have limited penetrating ability • α particles will only travel a few cm in air and will be stopped by thin cardboard • α particles have a +2 positive charge due to the 2 protons they contain Beta decay • Two types 1. β - occurs naturally 2. β + -produced in laboratories 1. β • β - decay results from a neutron in the nucleus transforming into a proton and an electron • The proton remains in the nucleus and the electron (β -) is emitted

• The decay equation for the β- decay of caesium-137 is as follows ;

2. Β + • β + decay is also called positron emission • A proton converts to a neutron and a positron (a positive electron) • The neutron remains in the nucleus and the positron (β +) is emitted

Ionising ability • Radiation particles can interact with the electrons of other atoms resulting in the formation of ions (charged atoms) • This is known as the “ionising ability” • Many of these interactions form unstable ions which can cause undesirable chemical reactions ! and / or cell damage • Rapidly dividing cell types are the most affected

•The decay equation for the β + decay of sodium-22 is as follows;

• β particles are very light compared to α particles and have speeds ranging from that of an α particle to almost the speed of light • They can travel a few metres in air but will be stopped by metal sheets as thin as 1 mm Gamma decay • A gamma ray, γ, is high-energy electromagnetic radiation that is emitted from the nuclei of radioactive atoms. • Gamma rays usually accompany an alpha or beta emission. • Gamma rays have a very high frequency, have no mass and travel at the speed of light (3.0 x 108 m s-1) • They have no electric charge. • Their high energy and uncharged nature makes them a very penetrating form of radiation. • Gamma rays can travel an almost unlimited distance through air and even a few centimeters of lead or a metre of concrete would not completely absorb a beam of gamma rays.

Half-life • The rate of decay of a radioisotope is measured by its half-life. • The half-life, t1/2, of a radioisotope is the time that it takes for half of the nuclei in a sample of the ! radioisotope to decay. • The “activity” of a sample indicates the number of radioactive decays that are occurring in the sample each second. • Activity is measured in becquerels (Bq) where 1 Bq = 1 disintegration per second.

• The activity of any radioactive sample will decrease with time. Over a half-life, the activity of a sample will halve....